Silane
From Wikipedia, the free encyclopedia
| Silane | |
|---|---|
 |
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| IUPAC name | Silane |
| Other names | Silicon tetrahydride Silicon hydride Monosilane Silicane |
| Identifiers | |
| CAS number | [7803-62-5] |
| UN number | 2203 |
| RTECS number | VV1400000 |
| Properties | |
| Molecular formula | SiH4 |
| Molar mass | 32.12 g mol−1 |
| Appearance | Colourless gas |
| Density | ? kg m−3 (solid) 0.7 g/ml (liquid) 1.342 g L−1 (gas) |
| Melting point |
88 K (−185°C) |
| Boiling point |
161 K (−112°C) |
| Solubility in water | Insoluble |
| Structure | |
| Molecular shape | tetrahedral |
| Dipole moment | 0 D |
| Thermochemistry | |
| Std enthalpy of formation ΔfH |
-1615 kJ mol−1 |
| Standard molar entropy S |
283 J mol−1 K−1 |
| Hazards | |
| Main hazards | low toxicity, avoid exposure to skin, irritant, may cause redness and swelling, extremely flammable, pyrophoric |
| NFPA 704 |
 4
1
3
|
| Flash point | N/A |
| Autoignition temperature |
294 K (21°C) |
| Related compounds | |
| Related hydrides | methane |
| Related compounds | disilene |
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references |
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Silane is a chemical compound with chemical formula SiH4. It is the silicon analogue of methane. At room temperature, silane is a gas, and is pyrophoric — it undergoes spontaneous combustion in air, without the need for external ignition.[1] However, the difficulties in explaining the available (often contradictory) combuistion data are ascribed to the fact that silane itself is stable and that the natural formation of larger silanes during production, as well as the sensitivity of combustion to impurities such as moisture and to the catalytic effects of container surfaces causes its pyrophoricity.[2][3] Above 420°C, silane decomposes into silicon and hydrogen; it can therefore be used in the chemical vapor deposition of silicon.
More generally, a silane is any silicon analogue of an alkane hydrocarbon. Silanes consist of a chain of silicon atoms covalently bound to hydrogen atoms. The general formula of a silane is SinH2n+2. Silanes tend to be less stable than their carbon analogues because the Si–Si bond has a strength slightly lower than the C–C bond. Oxygen decomposes silanes easily, because the silicon-oxygen bond is quite stable.
There exists a regular nomenclature for silanes. Each silane's name is the word silane preceded by a numerical prefix (di, tri, tetra, etc.) for the number of silicon atoms in the molecule. Thus Si2H6 is disilane, Si3H8 is trisilane, and so forth. There is no need for a prefix for one; SiH4 is simply silane. Silanes can also be named like any other inorganic compound; in this naming system, silane is named silicon tetrahydride. However, with longer silanes, this becomes cumbersome.
A cyclosilane is a silane in a ring, just as a cycloalkane is an alkane in a ring.
Branched silanes are possible. The radical ·SiH3 is termed silyl, ·Si2H5 is disilanyl, and so on. Trisilane with a silyl group attached to the middle silicon is named silyltrisilane. The nomenclature parallels that of alkyl radicals.
Silanes can also incorporate the same functional groups as alkanes, e.g. –OH to make a silanol. There is (at least in principle) a silicon analogue for all carbon alkanes.
Contents |
[edit] Production
Industrially, silane is produced from metallurgical grade silicon in a two-step process. In the first step, powdered silicon is reacted with hydrogen chloride at about 300°C to produce trichlorosilane, HSiCl3, along with hydrogen gas, according to the chemical equation:
- Si + 3HCl → HSiCl3 + H2
The trichlorosilane is then boiled on a resinous bed containing a catalyst which promotes its disproportionation to silane and silicon tetrachloride according to the chemical equation:
- 4HSiCl3 → SiH4 + 3SiCl4
The most commonly used catalysts for this process are metal halides, particularly aluminium chloride.
[edit] Properties
Silane has a repulsive smell.[4]
Silane has recently been shown to act as superconductor under extremely high pressures (96 and 120 GPa), with a transition temperature of 17 K.[5] Unfortunately, there was briefly an EE Times article that grossly exaggerated this achievement and claimed that room-temperature superconductivity had been achieved.
[edit] Applications
Several industrial and medical applications exist for silanes. For instance, silanes are used as coupling agents to adhere glass fibers to a polymer matrix, stabilizing the composite material. They can also be used to couple a bio-inert layer on a titanium implant. Other applications include water repellents, masonry protection, control of graffiti,[6] applying polycrystalline silicon layers on silicon wafers when manufacturing semiconductors, and sealants. Semiconductor industry alone used about 300 metric tons per year of silane in the late 1990s.[3] More recently, a growth in low-cost solar panel manufacturing has lead to substantial consumption of silane for depositing amorphous silicon on glass and other surfaces.
Silane is also used in supersonic combustion ramjets to initiate combustion in the compressed air stream.
Silane and similar compounds containing Si-H-bonds are used as reducing agents in organic and organometallic chemistry.[7]
"Mars sand" exposes regular sand to trimethylhydroxysilane vapors to make the sand waterproof.
Silane may be used to fabricate a super-compressed, superconducting compound.[5]
[edit] Safety and precautions
A number of fatal industrial accidents produced by detonation and combustion of leaked silane in air have been reported.[8][9][10] Dilute silane mixtures with inert gases such as nitrogen or argon are even more likely to ignite when leaked into open air, compared to pure silane: even a 1% mixture of silane in pure nitrogen easily ignites when exposed to air.[11] Unlike methane, silane is also fairly toxic: the lethal concentration in air for rats (LC50) is 0.96% over a 4-hour exposure. In addition, contact with eyes may form silicic acid with resultant irritation.[12]
[edit] References
- ^ Emeléus, H. J. and Stewart, K. (1935). "The oxidation of the silicon hydrides". Journal of the Chemical Society: 1182 - 1189. doi:.
- ^ Koda, S. (1992). "Kinetic Aspects of Oxidation and Combustion of Silane and Related Compounds". Progress in Energy and Combustion Science 18 (6): 513-528. doi:.
- ^ a b Timms, P. L. (1999). "The chemistry of volatile waste from silicon wafer processing". Journal of the Chemical Society - Dalton Transactions (6): 815-822. doi:.
- ^ CFC Startec properties of Silane
- ^ a b M. I. Eremets, I. A. Trojan, S. A. Medvedev, J. S. Tse, Y. Yao (2008). "Superconductivity in Hydrogen Dominant Materials: Silane". Science 319 (5869): 1506 - 1509. doi:.
- ^ Graffiti protection systems
- ^ Reductions of organic compounds using silanes
- ^ Chen, J. R. (2002). "Characteristics of fire and explosion in semiconductor fabrication processes". Process Safety Progress 21 (1): 19-25. doi:.
- ^ Chen, J. R.; Tsai, H. Y.; Chen, S. K.; Pan, H. R.; Hu, S. C.; Shen, C. C.; Kuan, C. M.; Lee, Y. C.; and Wu, C. C. (2006). "Analysis of a silane explosion in a photovoltaic fabrication plant". Process Safety Progress 25 (3): 237-244. doi:.
- ^ Chang, Y. Y.; Peng, D. J.; Wu, H. C.; Tsaur, C. C.; Shen, C. C.; Tsai, H. Y.; and Chen, J. R. (2007). "Revisiting of a silane explosion in a photovoltaic fabrication plant". Process Safety Progress 26 (2): 155-158. doi:.
- ^ Kondo, S.; Tokuhashi, K.; Nagai, H.; Iwasaka, M.; and Kaise, M. (1995). "Spontaneous Ignition Limits of Silane and Phosphine". Combustion and Flame 101 (1-2): 170-174. doi:.
- ^ See MSDS for silane.
